Stabilization and preservation of hydrogen ion concentration in both intra-cellular and extra-cellular body fluids is critical to the health of vertebrates. Considerable attention and scientific study have been devoted to this crucial aspect of maintaining internal homeostasis in humans.
Chemical substances which undergo reversible protonation serve as hydrogen ion buffers. Three chemical systems are primarily involved in vertebrates:
1) the CO.sub.2 -HCO.sub.3 balance; PA1 2) the protein zwitterion system, which is mutually dependent on the CO.sub.2 -HCO.sub.3 balance and which includes contributions made by the dissociation of the imidazole moiety of the histidine component of plasma proteins and hemoglobin; and PA1 3) the PHO.sub.3 /PO.sub.4 phosphate buffer system. Of these three chemical systems, the phosphate buffer system is the least effective, primarily because of the phosphate buffer system's inefficiency as a buffer above pH 7.2, and also because excessive phosphate causes undesirable side effects. For these reasons and others, phosphate buffers have had very limited clinical utility.
The use of hypothermia in cardiovascular surgery in the early 1950's necessitated an understanding of the relationship between body temperature and normal acid-base equilibria. One critical factor in the relationship between body temperature and pH involves dissociation constants (pK), which change with temperature. This change is substance-specific and is quite variable among different substances.
The change in pK of water is the prime determinant of the pH of neutrality in all living tissues because living tissue is mostly water. As tissue temperature decreases the dissociation of water therein decreases, the pH increases and the pH of neutrality rises. The term "pKa" describes the midpoint of the effective pH range of a given buffer composition.
Buffers are generally efficient only in narrow pH ranges which depend, in part, on their dissociation constants. For example, tris(hydroxymethyl) aminomethane, referred to herein as TRIS, is efficient in the pH range of 7.8 to 8.6, protein buffers are efficient in the range of 7.2 to 7.8 and bicarbonate is efficient in the range of 6.1 to 6.6.
In vertebrate tissue, the intra-cellular pH is regulated so that the hydroxyl ion concentration [OH.sup.- ] is equal to hydrogen ion concentration [H.sup.+ ], making the ratio of the concentration of OH- to H+ equal to 1. While a concentration ratio of 1 defines the pH of neutrality of water (pN), the pH of neutrality varies with temperature. Thus, pN at 37.degree. C. is about 6.8, at 27.degree. C. is about 6.95, and at 17.degree. C. is about 7.17. It is significant to note, however, that the blood and other fluids circulating in the body maintain a pH higher than pN by a factor k which is species specific. Thus, the pH of blood is pN+k. The cell membranes in humans regulate the k value to equal 0.6 pH units.
It is important to understand the relationship of cells and body fluids to pN at various temperatures to evaluate buffer efficiency because the pK of the buffer must rise as temperature falls so that the buffer will remain efficient throughout the range of the rising pH of cell and body fluids. In view of the importance of this relationship, many skilled in the art consider that the buffer pK should parallel pN at about 0.6 pH units higher to maintain efficiency. Unfortunately, the pK of the bicarbonate and phosphate buffers do not change proportionately to that of pN and, as cooling progresses, both buffers are in a pH range increasingly remote from the tissue pH range in which they are expected to work.
In contrast, the pK of imidazole, the buffer moiety of histidine almost superimposes the changing pN of water as temperature lowers. This explains the value of the histidine-rich plasma proteins and hemoglobin as buffers in human tissues and fluids, not only at homothermia but also at all viable hypothermic temperatures.
A convenient and useful designation concerning the buffering capacity of a buffer solution is the designation of the buffer's Van Slyke buffer value .beta.. The Van Slyke buffer value .beta. indicates the resistance of the buffer to change in pH upon addition of an acid or base, and is defined by the ratio .DELTA.B/.DELTA.pH, where B is the increment of completely dissociated base or acid in gram-equivalents per liter required to produce a unit change in pH. Thus, Slykes are the units of buffer strength (.beta.) derived form the formula .beta.=B/1pH, where B is millimoles of 1N acid or base. The non-bicarbonate buffer value of human blood is estimated at 28 Slykes of which 8 are due to plasma proteins and 20 are due to hemoglobin.
In healthy persons the acid-base equilibria of tissue is maintained largely by the interplay of the bicarbonate and protein buffer systems in accordance with the following relationships: ##STR1## Protein buffer systems, denoted herein as "Pr.sup.- ", include hemoglobin which serves as a reservoir of H.sup.+ ions, accepting or releasing H+ ions as equilibria dictate. In the event of hypothermia, bicarbonate becomes ineffective as a buffer and hemoglobin is the most effective buffering mechanism in the body.
A variety of clinical conditions threaten the acid-base balance. Acute acidosis can result from circulatory or respiratory insufficiencies or metabolic disorders. Circulatory insufficiency can be caused by coronary occlusion with myocardial hypoxia, traumatic wound of the heart, constrictive pericarditis or toxic myocarditis, hemorrhage, general body hypothermia, occlusive vascular disease, severe diarrhea and any severe trauma causing diminution in circulating blood volume. Respiratory insufficiency can result from thoracic wounds or injury, pulmonary inflammation, emphysema or a pulmonary embolism. Metabolic acidosis can result from diabetes, renal failure, inborn metabolic disorders or hepatic insufficiency.
Some acidotic conditions, for example severe diarrhea, can be efficiently treated by salt and water replacement. However most acidotic conditions cannot be so treated efficiently. One such condition is acidosis which universally accompanies cardiopulmonary bypass surgery with circulatory arrest in infants and adults. A safe, efficient buffer solution for use in cardiopulmonary bypass surgery has yet to be marketed.
During hypothermic cardiopulmonary bypass surgery it is desirable to dilute the blood significantly for rheological reasons. For instance, the heart-lung machine is primed with crystalloid solutions often containing blood to augment oxygen transport and thus to decrease perfusion rate. Recently, increasingly colder temperatures have been induced to improve myocardial tolerance to ischemia and to diminish body tissue oxygen requirements. In adults, as the crystalloid prime mixes with the patient's blood, the resulting perfusion fluid has an adequate oxygen transport capability but loses significant buffering capacity due to a diminished hemoglobin and plasma concentration. In infants when crystalloid prime is used without the addition of blood, the resulting fluid is an inadequate oxygen transport medium at temperatures above 15.degree. C. Under these conditions, hypoxia, plus lack of buffering capacity, leads to the lactic acidosis of hypoxia. As a result, acidosis is a fairly common occurrence in cardiac operating rooms especially because of the decreased use of blood bank blood due to the risk of AIDS. Since lack of buffer capacity contributes to lactic acidosis in hypoxia, a pump prime should contain a buffer even if the oxyhemoglobin content is lowered.
Since World War I NaHCO.sub.3 solution has increasingly become the only buffering agent available. With the vigorous methods of resuscitation currently in vogue using pre-loaded syringes, many patients with circulatory crises, severe trauma or other causes of lactic acidosis are receiving large doses of NaHCO.sub.3. In some cases, the large doses of NaHCO.sub.3 amount to an overload. The clinical evidence that such therapy is beneficial is extremely scanty; on the contrary, the current literature increasingly warns of the danger of such therapy. It has even been noted that no apparent clinical relationship exists between the amount of bicarbonate administered and eventual patient outcome.
In patients with respiratory problems, the failure to eliminate the CO.sub.2 may cause serious acidosis with dangerous levels of P.sub.CO2, the partial pressure of CO.sub.2 in the blood. When sodium bicarbonate is given intravenously, the CO.sub.2 generated diffuses more readily than bicarbonate across cell membranes and into the cerebrospinal fluid, thus potentially augmenting cerebral damage during hypoxic episodes. Hyperosmolality resulting from concentrated bicarbonate infusion can also cause myocardial arrhythmias, low capillary flow and decreased myocardial contractility.
The high sodium content of bicarbonate causes extracellular sodium retention when excessive amounts cannot be excreted. Excessive sodium retention results in excessive water retention, which in turn results in tissue edema and malfunction of various organs. Moreover, blood volume increases and oxygen transfer in impeded, thus increasing circulatory demand which increases the work load on the heart. In addition, the sodium overload can cause a delayed metabolic alkalosis which can lead to tetany and convulsions. In infants, this condition is thought to contribute to neonatal intracranial hemorrhage.
The absence of clinical evidence supporting the use of bicarbonate in the treatment of ketoacidosis, hypoxic lactic acidosis or circulatory arrest is not surprising in light of the buffering deficiencies of NaHCO.sub.3. Its pKa varies with temperature between 6.1 at 37.degree. C. and 6.2 at 17.degree. C. Thus, bicarbonate's effective buffer range of 5.7 to 6.8 lies almost totally outside the pH range of most clinical problems and its low .DELTA.pk/.degree.C. renders it ineffective in hypothermic states.
A recent series of papers, Arieff et al., "Pathophysiology of Experimental Lactic Acidosis in Dogs", Am J. Physiol., 1980, 239: F 135-142; Arieff et al., "Systemic Affects of NaCHO.sub.3 in Experimental Lactic Acidosis in Dogs", Am. J. Physiol. 1982, 242: F 586-591; Graf et al. "Metabolic Effects of Sodium Bicarbonate in Hypoxic Lactic Acidosis in Dogs", Am. J. Physiol., 1985, 249:F 630-635; and Graf et al. "Evidence for a Detrimental Effect of Bicarbonate Therapy in Hypoxic Lactic Acidosis", Chest, 1983, 83: 712-716, describe two animal models of lactic acidosis in which the efficiency of sodium bicarbonate therapy is examined. These experimental studies clearly indicate that under conditions of induced metabolic or hypoxic lactacidemia, intravenous administration of sodium bicarbonate worsens rather than alleviates the metabolic and hemodynamic consequences.
In an attempt to avoid the profound acute elevations in P.sub.CO2 induced by intravenous bicarbonate, a salt solution containing 0.33M bicarbonate and 0.33M carbonate has been proposed. While this solution, termed "carbicarb," diminishes or ablates the immediate rise in P.sub.CO2, after 15 minutes the carbicarb has the same effect achieved by infusing an equal amount of bicarbonate solution. Moreover, the sodium overload is the same and both solutions are ineffective to raise blood pH in experimental shock, lactic acidosis, cardiopulmonary arrest and certain other conditions. A recent study, Kette et al., "Buffer Solutions May Compromise Cardiac Resuscitation by Reducing Coronary Perfusion pressure", JAMA 1991, 266:2121-2126, presents evidence that the previously documented failure of infused bicarbonate and carbicarb to protect intramyocardial pH during porcine cardiac arrest may be due to depression in coronary blood flow caused by the extreme hyperosmolality of these two "buffers" at 2000 milliosmols per liter.
Clearly, the widely used bicarbonate solutions are therapeutically inadequate and are dangerous when used as alkalinizing agents. Unfortunately, there are no proven non-protein buffers currently available for use in clinical therapy.
Although blood and tissue byproduct buffers exist, such as protein buffer substitutes, they are prepared from blood or blood products of other species and therefore contain foreign protein which can cause a hyperimmune response. Moreover, such protein buffer substitutes are expensive and the disposal from the patient's body of degradation products can cause significant problems.
Synthetic buffering agents have also been investigated. One possible synthetic agent, TRIS, received extensive pharmacological investigation and clinical trials in the early 1960's. TRIS potentially can directly neutralize the weak acids of metabolic acidosis. However, TRIS is a strong base with a pH of 10, a pKa of 8.3 at 20.degree. C., and a .DELTA.pK/.degree.C. of about 0.020. The intravenous administration of TRIS causes severe vasospasm, phlebitis and local cellular necrosis. Neutralization with HCl to avoid these complications nullifies TRIS' capacity to buffer carbonic acid and to generate bicarbonate. There is little current clinical, theoretical or experimental evidence supporting a significant therapeutic role for TRIS. Essentially, TRIS has a pKa too high to be effective in the pH range of clinical acidosis. However, because of its high alkalinity and absence of sodium, TRIS might serve as a useful component in a compound with other synthetic buffers.
A family of N-substituted aminosulfonic acids have been synthesized and shown to be potentially useful buffering agents. However, it is believed that none of these synthetic buffering agents have had whole animal toxicity or pharmacological studies performed. It is further believed that these synthetic buffering agents have been used only in tissue culture, virus, enzyme and thermodynamic studies.
In 1966, Good and his associates introduced ten new hydrogen ion buffers, four of which were synthetic zwitterion aliphatic amines derived from taurine or glycine using bromoethanesulfonate. See Good et al., "Hydrogen Ion Buffers for Biological Research," Biochem., 1966, 5:467-477. Later, Ferguson and his colleagues introduced five similar new buffers synthesized from a different source. See Ferguson et al. "Hydrogen Ion Buffers for Biological Research," Anal Biochem, 1980, 104:300-310. These aminosulfonic acid buffers together with other synthetic compounds form a family of about sixteen synthetic buffers each with a known pKa, .DELTA.pK.degree./C. and neutralization equivalent. A few of these synthetic buffers have been used successfully in laboratories as precise buffers in studies involving enzyme dynamics and virus, bacterial and cell tissue cultures.
As early as 1971, the importance of differences in the pKa of various members of the aminosulfonic acid family in designing tissue culture media with optimal buffering capabilities at different pH levels has been recognized. For example, Eagle H., Science, 1971 174:500-503 showed that three different aminosulfonic acids improved the buffering stability of a 24 millimolar solution of sodium bicarbonate, with sodium phosphate sometimes supplemented. However, of the 16 biological buffers evaluated, 8 were too toxic to use for cell culture in concentrations as low as 0.020M. The other 8 buffers showed significant suppression of growth of certain cell lines. The fact that some of the aminosulfonic acids were cell toxic in extremely low concentration posed a challenge to investigation in the whole animal field of the current invention.
Combinations of two aminosulfonic acids, N-(2-hydroxyethyl)piperazine-N'-(2-ethanesulfonic acid) and 3-(N-morpholino) propanesulfonic acid and their respective sodium salts have also been used to produce very precise pH solutions for use in standardizing pH instruments. See Eur. Pat. Appl RP 341,793, IT Appl. 88/20,560, 13 May 1988, to Manzoni, A. and Belluati, M. This patent discloses aminosulfonic acid in combination with its sodium salt to create a buffer solution with a precise pH but no consideration is given to the effect of temperature on pH nor to any of the biological characteristics needed in a clinical buffer.
Thus, although lactic acidosis is a very common and often fatal clinical disorder, not a single effective therapeutic agent is currently available to prevent or counteract it. Furthermore a buffer composition with a pKa about 7.2 at 25.degree. C. is urgently needed to counteract the myocardial acidosis resulting from hypothermic heart surgery.
While approximately 16 synthetic compounds having buffering capability are commercially available, such synthetic compounds have been used extensively only in laboratory studies as components in buffer compositions to stabilize pH in culture media for cell, bacterial or viral studies. The effect of temperature on the pKa of buffer systems has been completely unaddressed.
It is believed that no attempt has been made to evaluate synthetic buffers in vertebrates, especially humans. It is currently not known whether any members of the family of N-substituted aminosulfonic acids and/or their sodium salts are non-toxic when rapidly administered intravenously to laboratory test animals in large doses. It is not known if compositions of aminosulfonic acids and their salts can be devised to have pH, pKa, .DELTA.pk/.degree.C., osmolality and sodium ion concentration characteristics to treat different forms of clinical hypoxic acidosis specifically and effectively.
It is the principal object of the present invention to produce effective synthetic therapeutic agents suitable for use as buffering agents in vertebrates. It is a further object of the present invention to produce buffering agents suitable for use in vertebrates which can counteract lactic acidosis in vertebrates. It is yet a further object of the present invention to produce buffering agents especially useful for counteracting the myocardial acidosis of hypothermic heart surgery and the generalized lactic acidosis of hemorrhage and traumatic shock.
It is a still further object of the present invention to provide methods for administering a synthetic buffering agent in large doses to counteract lactic acidosis in vertebrates. It is a still further object of the present invention to provide methods for devising synthetic buffering agents suitable for use in counteracting lactic acidosis in vertebrates and having appropriate pH, pKa, .DELTA.pK/.degree.C., osmolality and sodium ion characteristics designed to treat different forms of clinical hypoxic acidosis specifically and effectively.